Electron Configuration Notes
Quantum Numbers
 Quantum numbers are an address system for electrons.
 Each electron has four quantum numbers that give it a specific address within the
atom.
 The quantum numbers are represented by the letters n, l, m and ms.
Principal Quantum Number (n)
 The principal quantum number, designated “n,” indicates the average distance
from the nucleus where an electron resides—or the principle energy level in
which it resides.
 The set of orbitals within the same n is called an electron shell.
 An electron’s “n” is determined by locating the row of the periodic table in which
the element is located.
 Example: Potassium is on row 4 of the periodic table, thus having a principal
quantum number (n) of 4.
Azimuthal Quantum Number (l)
 The azimuthal quantum number, l (a lowercase letter L), describes the sublevels
of each principle energy level.
 The sublevels are symmetrical shapes that surround the nucleus.
 Each sublevel has a unique shape represented by the letters s, p, d, and f.
 The sublevels are assigned the following numerical values for l:
s=0
p=1
d=2
f=3
 The set of orbitals that have the same n and l values is called a subshell.
 Each subshell is written with a number (n value) and a letter (l value) of s, p, d, or
f.
 Example: the orbitals that have n=4 and l =2 are called “4d orbitals” and are in the
4d subshell.
You can find an atom’s l by identifying which “block” of the periodic table it is in.
Magnetic Quantum Number (m)
 Describes the spatial orientation of the orbitals within an atom
 For each l, there are 2l + 1 orbitals
◦ Example:
For l =2, there are 2(2) + 1 = 5 orbitals
 “m” can be anywhere from –l to +l.
◦ Example:
For l =2, m can be: -2, -1, 0, +1, +2
Electron Spin Quantum Number (ms)
 Each orbital holds only two electrons.
 These two electrons have opposite spin.
 Values for the ms are either +1/2 or -1/2
Orbital Types (a.k.a. Energy Sublevels)
 s orbital:
◦ Simplest of all orbitals
◦ Has a spherical shape
◦ Unlike other sublevels, the s sublevel contains only one orbital. It can hold
2 electrons.
 p orbitals:
◦ Contains three dumbbell shaped orbitals arranged around the x, y, and z
axes.
◦ Each shell of n=2 or greater has three p orbitals

d orbitals:
◦ There are five orbitals in the d sublevel.
◦ Any shell of n=3 or greater has five d orbitals
◦ Since each orbital holds two electrons, there can be a maximum of ten
electrons in each d sublevel.
 f sublevel:
◦ There are seven orbitals in the f sublevel.
◦ Since each orbital can hold two electrons, each f sublevel can hold a
maximum of 14 electrons.
◦ Only ground state atoms with n=4 or greater have an f sublevel.
◦ Very complicated orbital shape—much more complicated than d orbitals.
Electron Configuration
 The way an atom’s electron’s are distributed among its orbitals is called its
“electron configuration.”
 The electrons want to be in the lowest possible energy level, but they can’t all
crowd into the 1s orbital.
 The Pauli Exclusion Principle tells us that each orbital can hold, at most, 2
electrons.
 Electrons fill orbitals in order of increasing orbital energy.
 Diagonal Rule: helpful hint to remember the order in which orbitals are filled.
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
5f
6f
 Practicing electron configuration:
◦ Lithium:
 Lithium has 3 electrons (element 3 on periodic table).
 The orbital with lowest possible energy, the 1s, is filled first and it
holds 2 electrons.
 The remaining electron must fill the next available orbital, the 2s.
Li: __ __ __ __ __ __ __
1s
2s
2p

Lithium has an electron configuration of 1s22s1
 More practice with electron configuration:
Electron configuration of Boron:
Boron has 5 electrons (element 5 on the periodic table).
B: ____
1s
____
2s
___ ___ ___
2p
Summary of Boron’s electron configuration: 1s22s22p1
 More practice….
◦ Nitrogen has seven electrons (element 7 on periodic table).
N: ____ _____ ___ ___ ___
1s
2s
2p
Summary: 1s22s22p3
**Notice, that the electrons are unpaired in the 2p orbitals. Hund’s Rule says that for
like orbitals, the lowest energy is achieved when the number of electrons with the
same spin is greatest.
In other words, don’t pair electrons until you have to!
 Practice using Hund’s Rule:
◦ Electron configuration for oxygen:
◦ Oxygen has 8 electrons (element 8 on periodic table).
O: ____ ____ ____ ____ ____
1s
2s
2p
Summary: 1s22s22p4
 Silicon (element 14 on periodic table):
Si: ____ ____ ____ ____ ____ ____ ____ ____ ____
1s
2s
3s
2p
3p
Summary form Si: 1s22s22p63s23p2
We can also do Si another way:
Condensed Electron Configurations
We can use a short-hand way of showing Si’s electron configuration….
 The 2p subshell is filled at neon; neon has a stable electron configuration, with
eight electrons in its outermost shell forming an “octet”
 When 8 electrons fill the outermost shell, the atom is in a very stable
configuration.
 To use condensed electron configuration, we use the nearest noble-gas element
that has a lower atomic number. (To find our which noble gas to use, look at the
last column and use the noble gas right above the one on the same row as your
element)
 For example: Lithium would be Li: [He]2s1
Oxygen would be O: [He]2s22p5.
Silicon would be Si: [Ne]3s23p2